Foundations and Applications

Solution Chemistry

Photo by: Digital_Zombie

The majority of chemical processes are reactions that occur in solution.
Important industrial processes often utilize solution chemistry.
"Life" is the sum of a series of complex processes occurring
in solution. Air, tap water, tincture of iodine, beverages, and household
ammonia are common examples of solutions.

A solution is a homogenous mixture of substances with variable
composition. The substance present in the major proportion is called the
solvent, whereas the substance present in the minor proportion is called
the solute. It is possible to have solutions composed of several solutes.
The process of a solute dissolving in a solute is called dissolution.

Many common mixtures (like concrete) are
heterogeneous
—the components and properties of such mixtures are not distributed
uniformly throughout their structures. Conversely, solutions are said to
be
homogeneous
because they have uniform composition and properties. Solutions are
intimate and random homogeneous mixtures of atomic-size chemical species,
ions, or molecules.

In addition to their observed homogeneity, true solutions also have
certain other characteristics. For example, components of a solution never
separate spontaneously, even when a significant density difference exists
between the components. Solutions also pass through the finest filters
unchanged.

Droplets of a solution of water and oil, exposed to polarized light
and magnified.

The components of a solution distribute themselves in a completely random
manner, given sufficient time. For example, a lump of sugar dropped into a
glass of water dissolves, and eventually molecules of sugar can be found
randomly distributed throughout the water, even though no mechanical
stirring has been employed. This phenomenon, called diffusion, is similar
to the process of diffusion that occurs with gases. The molecules of sugar
(as well as those of water) must be in constant motion in the solution. In
the case of liquid solutions, the sugar molecules do not move very far
before they encounter other molecules; diffusion in a liquid is therefore
less rapid than diffusion in a gas.

Kinds of Solutions

Many commonly encountered solutions are those involving a solid that has
been dissolved in a liquid, but there are as many types of solutions as
there are different combinations of solids, liquids, and gases. Solutions
in which the solvent is a liquid and the solute is a gas, liquid, or solid
are very common. The atmosphere is a good example of a solution in which a
gaseous solvent (nitrogen) dissolves other gases (such as oxygen, carbon
dioxide, water vapor, and neon). Solutions of solids in solids are another
example, and these are encountered most often among the various metal
alloys.

Of all the liquid solvents used in the laboratory, in industry, and in the
home, water is the most commonly employed and is the best of the inorganic
solvents. The alcohols and numerous other types of compounds are
classified as organic solvents; many of these are used in dry cleaning
chemicals, nail polish removers, paint thinners, and many other similar
purposes.

Concentration

The concentration of a solution is defined as the amount of
solute
present in a given quantity of solvent. Very often scientists speak of
concentrated solutions, dilute solutions, or very dilute solutions, but
these designations give only a rough relative qualitative idea of
concentration. For example, a
"concentrated solution" contains a considerable quantity of
solute as compared with a "dilute solution." Although such
designations are only qualitatively useful, they are nevertheless widely
used.

The most common way to express concentration is on the basis of the weight
of solute per unit weight of solvent. For example, a salt solution may be
prepared by dissolving 1.64 grams of sodium chloride in 100 grams of
water. The concentration of this solution could also be expressed as
0.0164 grams of NaCl per 1 gram of water, or as 16.4 grams of NaCl per
1,000 grams of water. Thus, a statement of the concentration of a solution
does not imply anything concerning the amount of solute or the amount of
solvent present, but rather gives the ratio of solute to solvent in terms
of some convenient (and arbitrary) units. Because the weight of a sample
of a liquid is usually more difficult to determine experimentally than its
volume, a practical unit of concentration is the weight of solute in a
given volume of the solution; for example, a sugar solution may contain 50
grams of sugar per 100 milliliters of solvent.

Solubility

Solubility is a measure of the maximum amount of solute that can be
dissolved in a given amount of solvent to form a stable solution. The
composition of many solutions cannot be varied continuously because there
are certain fixed limits imposed by the nature of the substances involved.
Solid salt and sugar can be mixed in any desired proportions, but
unlimited quantities of sugar (or salt) cannot be dissolved in a given
amount of water; however, up to the solubility limit, solutions can be
produced in any desired proportion.

When the solvent contains a maximum quantity of solute, the resulting
solution is said to be saturated. The saturation point varies according to
the solute. For example, 100 grams of pure water at 25°C
(77°F) can dissolve no more than 35.92 grams of NaCl to form a
stable saturated solution, but this same amount of water at 25°C
dissolves only 0.0013 grams of calcium carbonate. The solubility in these
examples is expressed in grams of solute per 100 grams of water, but any
suitable units could be used. Water can dissolve any amount of a solute
less than that required for a saturated solution. Tables of the
solubilities of many substances can be found in various chemistry texts.

In some cases there is no upper limit to the amount of a solute that a
given quantity of solvent can dissolve, and these substances are said to
be miscible in all proportions. Completely miscible substances give
homogeneous mixtures (solutions); for example, a mixture of any two
gaseous substances is homogeneous. Often, liquids such as alcohol and
water can be mixed in all proportions to give homogeneous mixtures.

When a saturated solution has been achieved, a dynamic equilibrium exists
between the solute in solution and any undissolved solute. Molecules of
the solute (or atoms or ions, depending upon the nature of the solute) are
continuously going into solution, but since the solution is already
saturated, an equal number of molecules of the solute leave the solution
and redeposit on the excess solid solute. A state of equilibrium exists
when these two processes occur at the same rate, the net result being a
constant amount
of solute in solution. A saturated solution can therefore be defined more
precisely as a solution that is in equilibrium with an excess of the
solute at a given temperature.

In some instances it is possible to prepare a true solution that contains
an excess of the equilibrium amount of solute; this condition is called
supersaturation. Supersaturated solutions are unstable. If left
undisturbed, they may remain in this state for an indefinite period of
time. However, the excess solute can be brought out of solution by a
slight agitation or by the addition of any solid particle (dust, a small
crystal of solute, etc.) that can act as a center for crystal growth,
returning the solution to its normal saturated state.

Conditions That Affect Solubility

In general, three major factors—pressure, temperature, and the
nature of the solute and solvent—influence the solubility of a
solute in a solvent. Not all these factors are equally important in a
specific instance.

Pressure.
Changes in pressure have little effect on the solubility of solid or
liquid solutes in a liquid solvent, but pressure has a much greater
influence on the solubility of a gaseous solute. A commonly observed
phenomenon that supports this is the effervescence that occurs when the
cap of a bottle of ordinary soda water is removed. Soda water contains
carbon dioxide gas dissolved in water under pressure; when the cap is
removed, the pressure of the gas on the liquid is decreased to atmospheric
pressure. Since carbon dioxide gas leaves the solution at this lower
pressure, it follows that the solubility of carbon dioxide in water is
dependent upon the pressure of the carbon dioxide above the liquid. The
results of this simple observation are summarized in Henry's Law,
which states that at any specified temperature, the extent to which a gas
dissolves in a liquid is directly dependent upon the pressure of the gas.

Temperature.
In general, a change in temperature affects the solubility of gaseous
solutes differently than it does the solubility of solid solutes, because
the solubility of a gas in a liquid solvent decreases with increasing
temperature. With relatively few exceptions, the solubility of solids in
liquids increases with an increase in temperature. In some instances, the
increase in solubility is very large; for example, the solubility of
potassium nitrate in water at 25°C is about 31 grams of KNO
3
per 100 grams of water and about 83 grams of KNO
3
per 100 grams of water at 50°C (122°F). On the other hand,
the solubility of some solutes, such as ordinary table salt, shows very
little dependence on temperature. Often this difference in solubility can
be used as an advantage in the preparation, isolation, or purification of
substances by the process of crystallization. In general, it is not
possible to arrive at any reliable generalization concerning the influence
of temperature upon the solubility of liquids in liquids. In some cases
the solubility increases with an increase in temperature, in some cases it
decreases, and in others there is very little effect.

The nature of solute and solvent.
Crystalline substances consist of a regular arrangement of atoms,
molecules, or ions; in the latter case, the forces that hold the crystal
together are electrostatic in nature. For an ionic crystal to dissolve in
water, the water molecules must be able to shield the charges
of the positive and negative ions from each other. The attractive forces
between the ions in solution are less than those in the solid state
because of the solvent molecules; hence, the ions behave more or less
independently in solution. In general, the relative solubilities of ionic
substances are a measure of the magnitude of the electrostatic forces that
hold the crystals together.

Properties of Solutions

Pure liquids have a set of characteristic physical properties (melting
point, vapor pressure at a given temperature, etc.). Solutions in a
solvent exhibit these same properties, but the values differ from those of
the pure solvent because of the presence of the solute. Moreover, the
change observed in these properties in going from the pure solvent to a
solution is dependent only upon the number of solute molecules; these
properties are called colligative properties. The properties of a solvent
that show a predictable change upon the addition of a solute are melting
point, boiling point, vapor pressure, and osmotic pressure.

Melting and boiling points.
Solutions exhibit higher boiling points and lower melting points than the
parent solvent. The increase in boiling point and decrease in melting
point is dependent upon the number of solute particles in the solution.
The greater the number of solute particles (i.e., the concentration), the
greater the boiling point elevation and melting point depression. A common
application of this effect in some parts of the world is in the use of
antifreeze solutions in the cooling systems of automobiles in cold
climates. "Antifreeze" compounds are usually organic liquids
that are miscible with water so that large freezing point effects can be
attained.

Vapor pressure
All liquids exhibit a vapor pressure, the magnitude of which depends on
the temperature of the liquid. For example, water boils at 100°C,
which means that at 100°C the vapor pressure of water is equal to
the atmospheric pressure allowing bubbles of gaseous water (steam) to
escape from the liquid state. However, the vapor pressure of a solution
(at any temperature) is less than that of the solvent. Thus, boiling water
ceases to boil upon the addition of salt because the salt solution has a
lower vapor pressure than pure water. The salt solution will eventually
boil when the temperature of the solution increases bringing about an
increase in vapor pressure sufficient to again form bubbles. Note in this
example that the boiling point of water increases with the addition of
salt; thus, the boiling point elevation and the vapor pressure depression
are related.

Osmosis.
This property of solutions is perhaps the least familiar of the
colligative properties, but in a sense it is more important than those
already mentioned. In 1748 French clergyman and physicist Jean-Antoine
Nollet observed that certain animal membranes are selectively permeable to
different molecules. Since then, many examples of semipermeable membranes
have been discovered, including animal bladder or gut tissues, eggshell
lining, and certain vegetable tissues. A semipermeable membrane may be
defined as a material that allows molecules of one kind to pass through it
but prevents the passage of other kinds of molecules or allows the passage
of different kinds of molecules at different rates. Membranes often permit
the passage of solvent molecules and prevent the passage of solute
molecules.
The phenomenon of osmosis is of far-reaching importance in biology,
medicine, and related areas. Many animal and vegetable membranes are
semi-permeable, and the process of osmosis plays an important role in the
transfer of molecules through cell walls in biological processes. Osmosis
is responsible in part for the germination of seeds and for the rising of
sap into the branches and leaves of trees. The preservative action of
sugar solutions (e.g., preserves, jellies) is believed to depend upon
osmotic processes; bacteria are literally dehydrated.

I request your valuable reply. What I understood was when vapour pressure equels the atmospheric pressure solvent or solution boils. According to your article, solution has low vapour pressure. Then why the solution does not boil faster than solvent?

my name is trey and i am currently attenting monty tech high school. i dont really understand the difference of a solution and a mixture and also the difference of a solvent and a solute. i need to come up with the names of 10 different solutions so i can define them and make a list of what they are composed of. if you could help i would really apprichiate it. thank you for your time.

"The gas molecular are always in random motion. the molecules bump against each other and against
each other and against the walls of the container. The molecule exert a force against the wall
during collision, the force per unit area of the wall is called pressure. eg. say for instance
baloon.